Acids

1
Acids Bases
2
Properties of Acids

• Sour taste
• Change color of acid-base indicators (red in pH
  paper)
• Some react with active metals to produce hydrogen
  gas
• Ba(s) H2SO4(aq) BaSO4(s) H2(g)
• Some react with bases to neutralize and form salt
  and water
• H2SO4 (aq) 2NaOH(aq) Na2SO4 (aq)
  2H2O(l)
• Some are electrolytes

3
Examples of Acids

• Lemons and oranges - citric acid
• Vinegar - 5 by mass acetic acid
• Pop and fertilizer - phosphoric acid

4
Properties of Bases

• Bitter taste
• Change color of acid-base indicators (blue in pH
  paper)
• Dilute aqueous solutions feel slippery
• Ex. Soap
• Some react with acids to neutralize and form
  salt and water
• Some are electrolytes

5
Examples of Bases

• Soap - NaOH
• Household cleaners - NH3
• Antacids - Ca(OH)2, Mg(OH)2

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Arrhenius Acids

• Acids that increase the concentration of
  hydronium (H3O) in aqueous solutions
• HNO3(aq) H2O(l) H3O(aq)
  NO3-(aq)
• H NO3- H2O

acid
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Why do acids produce H3O?

• H is extremely attracted to the unshared pair of
  electrons on the water molecule so it donates
  itself to this molecule where it becomes
  covalently bonded. The ion formed is known as
  the hydronium ion (H3O)

H
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Arrenius Bases

• Bases that increase the concentration of
  hydroxide ions (OH-) in aqueous solutions
• NaOH(s) Na(aq) OH-(aq)

H2O
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Strength of Acids Bases

• Strong acids bases completely ionize in aqueous
  solutions
• H2SO4 H2O H3O HSO4-
• NaOH Na OH-
• Strong acids bases are strong electrolytes
• A list of strong acids bases can be found on
  pg. 460-461

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• Weak acids bases only partially break down into
  ions when in aqueous solutions
• HCN H2O H3O CN-
• NH3 H2O NH4 OH-
• Weak acids bases are weak electrolytes
• A list of weak acids bases can be found on pg.
  460-461

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Why can we drink H2O?

• Water self ionizes to form equal concentrations
  of H3O and OH-
• H2O(l) H2O(l) H3O(aq) OH-(aq)
• A substance is considered neutral when H3O
  OH-
• H3O concentration 1.0 x 10-7M
• OH- concentration 1.0 x 10-7 M

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When H3O OH-

• If H3O gt 1.0 x 10-7 M, the solution is acidic
• If OH- gt 1.0 x 10-7 M, the solution is basic
• To find the concentration of H3O or OH- in
  acidic or basic solutions, the following equation
  can be used
• 1.0 x 10-14 M2 H3O OH-
• 1.0 x 10-14 M2 ionization constant for H2O (Kw)

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Sample Problem

• A 1.0 x 10-4 M solution on HNO3 has been prepared
  for laboratory use.
• a. Calculate the H3O of this solution
• b. Calculate the OH- of this solution
• c. Is this solution acidic or basic? Why?
• d. Substitute H2SO4 as the acid. How would
  the calculations change?

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Sample Problem

• An aqueous 3.8 x 10-3 M NaOH solution has been
  prepared for laboratory use.
• a. Calculate the H3O of this solution
• b. Calculate the OH- of this solution
• c. Is this solution acidic or basic?
• Why?
• d. Substitute Ca(OH)2 as the base. How
  would the calculations change?

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Practice Problems

• Complete practice problems on pg. 484 1-4

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The pH scale

• The pH scale measures the power of the hydronium
  ion H3O in a solution
• The scale typically goes from 1-14 (although it
  can extend below or above it under extreme
  conditions)
• The following equations can be used to determine
  the pH or H3O of a solution
• pH -log H3O H3O antilog (-pH)
• H3O 1 x 10-pH

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pH gt 7 basic pH 7 neutral pH lt 7 acidic
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The pOH scale

• The pOH scale measures the power of the hydroxide
  ion OH- in a solution
• The scale typically goes from 1-14 (although it
  can extend below or above it under extreme
  conditions)
• The following equations can be used to determine
  the pOH or OH- of a solution
• pOH -log OH- OH- antilog (-pOH)
• OH- 1 x 10-pOH

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• pH pOH 14

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Sample Problems

• Calculate the pH of each of the following.
  Classify as acidic or basic.
• 1.3 x 10-5 M NaOH
• 1.0 x 10-4 M HCl

21
Sample Problems

• What is the H3O for each of the following?
  Classify as acidic or basic.
• pH 5.8
• b. pOH 8.9

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Sample Problems

• What is the OH- for each of the following?
  Classify as acidic or basic.
• H3O 9.5 x 10-10 M
• pOH 1.3

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Practice Problems

• Complete practice problems on
• pg. 487 1
• pg. 488 1-4
• pg. 490 1-4

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Expansion of the Acid-Base Theory

• Substances can still act as an acid or base if
  they are not dissolved in water to make a solution

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Bronsted-Lowry Acids

• A molecule or ion that is a proton (H) donor
• HCl(g) NH3(g) NH4(g) Cl-(g)

H donor
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Bronsted-Lowry Bases

• A molecule or ion that is a proton (H) acceptor
• HCl(g) NH3(g) NH4(g) Cl-(g)
• In a Bronsted-Lowry acid-base reaction, protons
  (H) are transferred from one reactant (the acid)
  another (the base)

H acceptor
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Monoprotic versus Polyprotic Acids

• Monoprotic acids can only donate 1 proton per
  molecule
• HCl(g) H2O(l) H3O(aq) Cl-(aq)

Monoprotic
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• Polyprotic acids can donate more than one proton
  per molecule
• H2SO4(aq) H2O(l) H3O(aq) HSO4-(aq)

Polyprotic
HSO4-(aq) H2O(l) H3O(aq) SO4-2(aq)
One additional proton can still be donated
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Conjugate acids bases

• A conjugate acid is the species that is formed
  when a Bronsted-Lowry base gains a proton
• A conjugate base is the species that remains
  after a Bronsted-Lowry acid has given up a proton
• HF(aq) H2O(l) F-(aq) H3O(aq)

acid
base
Conjugate acid
Conjugate base
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More examples

• CH3COOH(aq) H2O(l) H3O(aq)
  CH3COO-(aq)
• HCl(aq) H2O(l) H3O(aq)
  Cl-(aq)

CB
CA
acid
base
acid
base
CA
CB
Proton transfer reactions favor the production of
the weaker acid and base. Use table 15-6 on pg.
471 in your text to compare the relative
strengths of acids and bases
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Is H2O an acid or a base?

• H2O is amphoteric, it can react as either an acid
  or a base
• If H2O reacts with a compound that is a stronger
  acid than itself, it acts as a base
• If H2O reacts with a weaker acid, it will act as
  the acid
• H2SO4(aq) H2O(l) H3O(aq) HSO4-(aq)

Base H acceptor
NH3(aq) H2O(l) NH4(aq) OH-(aq)
Acid H donor
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OH- in a molecule

• When an OH- group is covalently bonded in a
  molecule, it is referred to as a hydroxyl group
• Hydroxyl groups are present in many organic
  compounds
• Ex. Acetic acid (HC2H3O2) or CH3COOH

Hydroxyl group
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How does the OH- make something acidic?

• In order for a compound with an OH- group to be
  acidic, H2O must be able to attract the H atom
  from the OH- group and act as a proton donor
• CH3COOH(aq) H2O(l) H3O(aq)
  CH3COO-(aq)

The more O atoms bonded to the OH- group, the
more acidic the compound is likely to be. Oxygen
is highly electronegative and will attract
electrons closer to it, making the OH- bond more
polar. This will allow H2O to steal the H
atoms more easily.
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Why are substances with OH- covalently bonded to
it sometimes not acidic?

• Ex. Acetic acid (CH3COOH) versus ethanol
  (C2H5OH)

Ethanol
Acetic acid
Acetic acid- the 2 O atom on the C atom draws
electron density away from the OH- group, making
the bond more polar. This allows the H to be
donated more easily Ethanol- this compound is
essentially neutral. It does not have a second O
atom to make the bond as polar. It would be
classified as a very weak acid because it is
harder to donate H.
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Further expansion of acid-base theory

• Substances can still act like an acid or base if
  they do not contain hydrogen at all

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Lewis acids bases

• A Lewis acid is an atom, ion, or molecule that
  accepts an electron pair to form a covalent bond
• Ag(aq) 2NH3(aq) H3N-- Ag--NH3
• A Lewis base is an atom, ion, or molecule that
  donates an electron pair to form a covalent bond

e- pair acceptor
e- pair donator
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Sample Lewis acid-base problem

• For the following equation, which reactant is the
  Lewis acid? Lewis base?
• BF3(aq) F-(aq) BF4- (aq)

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• BF3 is the Lewis acid because it is the e- pair
  acceptor
• F- is the Lewis base because it is the e- pair
  donor

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Review of acid-base categorization
Type Acid Base
Arrhenius H3O producer OH- producer
Bronsted-Lowry Proton (H) donor Proton (H) acceptor
Lewis e- pair acceptor e- pair donor
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Strong Acid-Base Neutralization

• When equal parts of acid and base are present,
  neutralization occurs where a salt and water are
  formed
• HCl(aq) NaOH(aq) NaCl(aq) H2O(l)

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Sample Problems

• H2CO3 Sr(OH)2
• HClO4 NaOH
• HBr Ba(OH)2
• NaHCO3 H2SO4

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Titrations

• When you have a solution with an unknown
  concentration, you can find it by reacting it
  completely with a solution of known concentration
• This process is known as titrating
• To perform a titration, an instrument called a
  buret can be used to precisely measure amounts of
  solution, drop by drop

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Titration Termonology

• Equivalence point - the point at which the known
  and unknown concentration solutions are present
  in chemically equivalent amounts
• moles of acid moles of base
• Indicator - a weak acid or base that is added to
  the solution with the unknown concentration
  before a titration so that it will change color
  or indicate when in a certain pH range (table
  16-6 on pg. 495 in your text will show various
  indicators and their color ranges)

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• End point - the point during a titration where an
  indicator changes color
• The 2 most common indicators we will use in our
  chemistry class will be
• Phenolphthalein - turns very pale pink at a pH of
  8-10
• Bromothymol blue - turns pale green at a pH of
  6.2-7.6

Phenolpthalein is clear at pHlt8, pale pink at pH
8-10 and magenta at pH gt10
Bromothymol blue
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Practice Titration for an unknown acid

• 1. Titrate 5.0 of mL of unknown HCl into a 250
  mL erlenmeyer flask - remember to document the
  starting amount and ending amount of acid on the
  buret to prevent error
• 2. Add 2 drops of indicator (phenolphthalein) to
  the flask - the color of the solution should be
  clear
• 3. Titrate with .5M NaOH, continuously swirling
  the flask, until the solution turns very pale
  pink for 30 seconds - remember to document the
  starting amount and ending amount of base on the
  buret
• 4. Mathematically determine the concentration of
  the unknown HCl solution by using the following
  equation

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Titration Equation

• MAVA MBVB
• MA molarity (mol/L) of acid
• VA volume in L of acid
• MB molarity (mol/L) of base
• VB volume in L of base
• molesA molesB
• 5. After calculating the molarity of the unknown
  acid experimentally, get the theoretical molarity
  and calculate error

48
Practice titration for an unknown base

• 1. Titrate 5.0 of mL of unknown NaOH into a 250
  mL erlenmeyer flask - remember to document the
  starting amount and ending amount of base on the
  buret to prevent error
• 2. Add 2 drops of indicator (phenolphthalein) to
  the flask - the color of the solution should be
  magenta
• 3. Titrate with .5M HCl, continuously swirling
  the flask, until the solution turns very pale
  pink for 30 seconds - remember to document the
  starting amount and ending amount of acid on the
  buret
• 4. Mathematically determine the concentration of
  the unknown NaOH solution by using MAVA MBVB
• 5. After calculating the molarity of the unknown
  base experimentally, get the theoretical molarity
  and calculate error

49
How do pH indicators work?

• Acid-base indicators are usually weak acids or
  bases that are in equilibrium and show color
  changes when a stress is applied
• HIn H In-
• In acidic solutions, the H concentration
  increases. The stress will cause a shift to the
  left (red color).
• In basic solutions, the OH- concentration
  increases. These ions will combine with H which
  will cause a shift to the right (blue color)

red
blue